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J. Chem. Eng. Data 2009, 54, 2028–2032
Physicochemical Properties of Aqueous Hydroxylamine Sulfate and Aqueous (Hydroxylamine Sulfate + Ammonium Sulfate) Solutions at Different Temperatures Sheng Fang,† Jian-Guo Zhang,‡ Jian-Qing Liu,‡ Ming Liu,† Sheng Luo,‡ and Chao-Hong He*,† Department of Chemical Engineering, Zhejiang University, Hangzhou 310027, China and Quzhou Jinhua Chemical Company, Limited, Quzhou 324004, China
The densities of an aqueous solution of pure hydroxylamine sulfate at T ) (283.15, 288.15, 293.15, 298.15, 303.15, 308.15, and 313.15) K and a solution of hydroxylamine sulfate containing ammonium sulfate at T ) (293.15, 298.15, 303.15, 308.15, and 313.15) K have been determined over a wide range of concentrations. The viscosities of binary and ternary mixtures of these compounds at T ) (298.15, 303.15, and 308.15) K are also presented. The Masson-type equation and extended Jones-Dole equation in terms of ionic strength are used to correlate the density and viscosity, respectively, of aqueous solutions of hydroxylamine sulfate. The ideal mixing rule of Young has been successfully applied to predict the ternary densities and viscosities of aqueous hydroxylamine sulfate + ammonium sulfate based on relevant binary solution properties.
Introduction Aqueous solutions of hydroxylamine sulfate in the presence of ammonium sulfate and sulfuric acid, produced by the Raschig process, are frequently used for polyamide manufacture.1 Pure hydroxylamine sulfate is also necessary in industry for manufacture of various chemicals, including herbicides and pharmaceuticals. Over the past century, various technologies have been proposed to separate hydroxylamine sulfate from the aqueous solution of ammonium sulfate produced by the Raschig process.1 The liquid ion-exchange extraction method has shown advantages because of its high efficiency and low cost and it can be performed in a continuous process.2,3 In the process, the Raschig mixtures are first naturalized with aqueous ammonia and then extracted by a nonpolar solvent containing organophosphate extractant.4 The industrial applications generally involve an organic phase containing organophosphate extractant and an aqueous phase containing hydroxylamine and ammonium salts. Moreover, the aqueous phase involves moderate to very high salt concentrations and different portions of hydroxylamine and ammonium salts in the extraction stages. Thus, reliable data and modeling on the physicochemical properties of the binary and ternary mixtures are needed over a wide range of concentrations and temperatures for simulation and optimization of the extraction process. Previous works have been devoted to the measurement of densities and viscosities of the organophosphate extractant in different nonpolar solvents.5,6 This work reports the properties of aqueous solutions of pure hydroxylamine sulfate and solutions of hydroxylamine sulfate + ammonium sulfate at different temperatures. The properties of aqueous hydroxylamine sulfate have been reported by Skvortsov et al.;7 however data for the industrially important ternary mixtures of hydroxylamine sulfate * Corresponding author. Phone: +86 571 87952709. Fax: +86 571 87951742. E-mail:
[email protected]. † Zhejiang University. ‡ Quzhou Jinhua Chemical Company, Limited.
+ ammonium sulfate + H2O have not been presented. For aqueous hydroxylamine sulfate solutions, the literature7 is quite old and not easy to obtain. In this work the solution densities of aqueous hydroxylamine sulfate were determined by highprecision vibrating-tube densimetry. The densities and viscosities of binary mixtures were correlated by the Masson-type equation8 and extended Jones-Dole equation9 in terms of ionic strength, respectively. We found that the linear additivity rule of Young10 can be successfully applied to predict the densities and viscosities of aqueous hydroxylamine sulfate + ammonium sulfate mixtures. The simple additivity rule assumes that the properties of a mixed solution at a given composition is additive in the properties of the individual binary solutions.
Experimental Section Materials. Hydroxylamine sulfate with a minimum purity in mass fraction of 0.995 was obtained from UBE (Japan). Ammonium sulfate with a minimum purity in mass fraction of 0.999 was obtained from the Juhua Group Corp. Polyamide Fiber Factory (Quzhou, China). The salt was placed and dried in a vacuum desiccator and used without further purification. Double-distilled deionized water was used for the preparation of aqueous solutions. Aqueous solutions were prepared from a known mass of salt and water in a stoppered glass bottle. The masses were recorded on a Sartorius Corp. BS 224S balance to a precision of ( 1 · 10-4 g. The liquid samples were filtered by using syringes equipped with 0.45 µm membrane filters before density and viscosity measurements were recorded. Experimental Measurements. The densities of the binary and ternary mixtures were measured with an Anton Paar DMA 4500 density meter whose measurement cell temperature was controlled automatically within ( 0.01 K of the selected value. It was calibrated at atmospheric pressure with double-distilled deionized water and dry air. Densities, both of water and of dry air, at the various working temperatures were given by the
10.1021/je8008649 CCC: $40.75 2009 American Chemical Society Published on Web 04/27/2009
Journal of Chemical & Engineering Data, Vol. 54, No. 7, 2009 2029 Table 1. Density of Hydroxylamine Sulfate in Water from T ) (283.15 to 313.15) K F/g · cm-3 m(SO42-)/mol · kg-1
I/mol · kg-1
T ) 283.15 K
T ) 288.15 K
T ) 293.15 K
T ) 298.15 K
0.0075 0.0380 0.0702 0.2199 0.3709 0.4314 0.5334 0.8069 0.9684 1.1647 1.4781 1.8505 2.4613 2.9977
0.0224 0.1141 0.2107 0.6596 1.1127 1.2942 1.6002 2.4206 2.9051 3.4940 4.4343 5.5516 7.3840 8.9932
1.00057 1.00384 1.00727 1.02247 1.03687 1.04264 1.05159 1.07532 1.08831 1.10323 1.12668 1.15194 1.18898 1.21805
0.99995 1.00318 1.00656 1.02157 1.03581 1.04122 1.05039 1.07390 1.08680 1.10158 1.12491 1.14999 1.18687 1.21584
0.99903 1.00223 1.00558 1.02043 1.03452 1.03997 1.04898 1.07196 1.08512 1.09981 1.12300 1.14792 1.18469 1.21355
0.99789 1.00103 1.00434 1.01905 1.03305 1.03836 1.04739 1.07020 1.08329 1.09789 1.12096 1.14549 1.18241 1.21121
T ) 303.15 K
T ) 308.15 K
T ) 313.15 K
0.0075 0.0380 0.0702 0.2199 0.3709 0.4314 0.5334 0.8069 0.9684 1.1647 1.4781 1.8505 2.4613 2.9977
0.0224 0.1141 0.2107 0.6596 1.1127 1.2942 1.6002 2.4206 2.9051 3.4940 4.4343 5.5516 7.3840 8.9932
0.99646 0.99959 1.00288 1.01748 1.03138 1.03673 1.04561 1.06830 1.08132 1.09585 1.11882 1.14234 1.18007 1.20879
0.99484 0.99795 1.00117 1.01570 1.02951 1.03477 1.04369 1.06630 1.07923 1.09370 1.11650 1.14116 1.17764 1.20632
0.99300 0.99609 0.99912 1.01304 1.02724 1.03285 1.04160 1.06415 1.07698 1.09142 1.11363 1.13836 1.17514 1.20399
Table 2. Polynomial Coefficients of Equation 5 for Densities of Aquous Hydroxylamine Sulfate Solution a1
a2
b1
b2
c1
c2
100 ARD
rmsd
0.012324
7.237611
0.002835
-1.979776
-0.000479
0.098221
0.014
0.00021
manufacturer in the instruction manual. The reproducibility of the density measurement was within ( 5 · 10-5 g · cm-3. The viscosities of the solutions were determined by using a Ubbelohde viscometer with capillaries (0.3 to 0.4) mm or (0.5 to 0.6) mm in diameter according to the viscosity of the samples. It was filled with experimental liquid and placed vertically in a DF-02 transparent-walled water bath (Fangao Scientific Co., Ltd., Nanjing, China), which was maintained constant to ( 0.01 K at the selected temperature. To minimize the evaporation loss during heating and measurement, the viscometer’s limb was closed using a rubber stopper with a small plastic tube in the center which was open to the atmosphere. An electronic digital stopwatch with a precision of ( 0.01 s was used for flow-time measurements. The viscosity of the solutions can be expressed as
η ) k1t - k2/t F
(1)
where F is the density of the liquid, t is the flow time, and k1, and k2 are the viscometer constants. The k1 and k2 parameters for both Ubbelohde viscometers were obtained by the measurements on double-distilled deionized water at (288.15, 293.15, 298.15, 303.15, and 308.15) K. It is convenient to get k1 ) 0.003486 and k2 ) -0.268447 for the (0.3 to 0.4) mm viscometer and k1 ) 0.009773 and k2 ) 0.810429 for the (0.5 to 0.6) mm viscometer by the linear regression of ηt/F with t2. Measurements were repeated at least four to five times for each solution and temperature. The reproducibility of the viscosity estimates was found to be within ( 0.1 %.
Results and Discussion Densities for (NH3OH)2SO4 + H2O Solutions. The densities of (NH3OH)2SO4 + H2O solutions are listed in Table 1. The
experimental data can be described by the Masson-type equation11,12 in terms of ionic strength in the form
F ) FH2O + aI + bI 3/2 + cI 2
(2)
with ionic strength in molality I defined as
I)
1 2
∑ mi zi2
(3)
i
where mi is the molality of ion i in moles per kilogram of pure solvent H2O and zi represents the ion charges. The temperature dependence of parameters a, b, and c in eq 2 will facilitate the density interpretation at other temperatures that are not determined by experiment. Thus, the following equation for the temperature dependence of the three parameters is used
P ) P1 +
P2 T
(4)
where P represents the parameters in eq 2, T is the absolute temperature in Kelvin, and P1 and P2 are constants. Thus, the temperature and concentration dependence of the density equation can be written as
(
F ) FH2O + a1 +
) (
)
(
)
a2 b2 3/2 c2 2 I + b1 + I + c1 + I T T T (5)
The values of the six parameters in eq 5 were estimated by the Levenberg-Marquardt algorithm with the help of SPSS software
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Journal of Chemical & Engineering Data, Vol. 54, No. 7, 2009
Table 3. Densities of Aqueous Hydroxylamine Sulfate + Ammonium Sulfate Mixture at Different Temperatures F/g · cm-3 T ) 293.15 K +
+
-1
m(NH3OH )/mol · kg
m(NH4 )/mol · kg
0.0000 0.5325 1.0693 1.3383 1.6073 2.1536 0.2535 1.0792 1.7861 3.1504 2.4698 1.7931 0.8918 1.2916 0.6971
a
-1
9.3100 8.7906 8.2826 8.0292 7.7681 7.2655 7.3433 6.5382 5.8686 0.2265 0.8970 1.5684 2.4525 0.8607 1.4479 rmsd 100 ARD
T ) 298.15 K a
T ) 303.15 K a
T ) 308.15 K a
T ) 313.15 K a
exp
calcd
exp
calcd
exp
calcd
exp
calcd
exp
calcda
1.21766 1.22288 1.22681 1.22841 1.23061 1.23487 1.19303 1.20033 1.20652 1.13488 1.12793 1.12116 1.11167 1.08424 1.07783
1.21644 1.22040 1.22459 1.22671 1.22871 1.23314 1.19229 1.19917 1.20535 1.13509 1.12816 1.12137 1.11207 1.08433 1.07780 0.00135 0.09
1.21579 1.22098 1.22487 1.22646 1.22868 1.23287 1.19106 1.19771 1.20446 1.13267 1.12566 1.11914 1.10980 1.08247 1.07610
1.21517 1.21908 1.22323 1.22532 1.22730 1.23168 1.19087 1.19766 1.20376 1.13302 1.12622 1.11956 1.11044 1.08263 1.07625 0.00095 0.06
1.21389 1.21904 1.22289 1.22446 1.22674 1.23082 1.18901 1.19625 1.20234 1.12977 1.12265 1.11585 1.10776 1.08055 1.07423
1.21366 1.21752 1.22163 1.22370 1.22565 1.22999 1.18922 1.19592 1.20194 1.13075 1.12407 1.11753 1.10857 1.08072 1.07447 0.00090 0.07
1.21194 1.21705 1.22088 1.22243 1.22475 1.22873 1.18683 1.19358 1.20009 1.12664 1.11880 1.11296 1.10549 1.07845 1.07169
1.21193 1.21575 1.21982 1.22186 1.22380 1.22809 1.18735 1.19396 1.19992 1.12828 1.12172 1.11530 1.10650 1.07860 1.07247 0.00103 0.08
1.20997 1.21504 1.21881 1.22033 1.22276 1.22660 1.18455 1.19173 1.19780 1.12339 1.11563 1.10967 1.10304 1.07585 1.06919
1.21001 1.21378 1.21780 1.21983 1.22174 1.22600 1.18528 1.19181 1.19770 1.12564 1.11919 1.11288 1.10422 1.07629 1.07029 0.00128 0.10
Calclulated by eq 8.
under the 0.95 confidence interval by minimizing the average relative deviation (ARD) as
ARD )
N |Fi,calcd - Fi,exp | 1 N i)1 Fi,exp
∑
{
N
(6)
}
∑ (Fexp - Fcalcd)2/N i
1/2
(7)
From the data listed in Table 2, it can be seen that the results show good agreement with the experimental values, which indicates that eq 5 can be applied to represent the densities within the concentration and temperature range. Densities for (NH3OH)2SO4 + (NH4)2SO4 + H2O Solutions. The densities of ternary mixtures for (NH3OH)2SO4 + (NH4)2SO4 + H2O are collected in Table 3. In the industrial extraction stage, different portions of hydroxylamine and ammonium salts exist in the aqueous phase. The densities reported in this work will be beneficial for the research of mass transfer during extraction, the simulation of the extraction process, and the design of extraction equipment. The simple additivity rule has been tested for its ability in the prediction of properties of the title system. The widely used Young’s rule10,12 can be applied as
F(IT) )
I(NH3OH)2SO4 IT
F(NH3OH)2SO4(IT) +
I(NH4)2SO4 IT
η/mPa · s
m(SO42-)
mol · kg-1 I/mol · kg-1 T ) 298.15 K T ) 303.15 K T ) 308.15 K
where N is the number of experimental points, Fi,exp represents the experimental densities, and Fi,calcd represents the densities calculated. The parameters a1, a2, b1, b2, c1, and c2 and rootmean-square deviations (rmsd) defined below are listed in Table 2.
rmsd )
Table 4. Viscosity of Hydroxylamine Sulfate in Water from T ) (298.15 to 308.15) K
F(NH4)2SO4(IT)
0.0204 0.0473 0.0702 0.2315 0.3709 0.4314 0.5334 0.8069 1.0436 1.3236 1.5892 1.9145 2.0882 2.8469
0.897 0.909 0.914 0.968 1.019 1.035 1.069 1.167 1.256 1.369 1.483 1.631 1.732 2.123
0.806 0.814 0.821 0.869 0.912 0.927 0.961 1.053 1.128 1.237 1.333 1.463 1.542 1.887
0.727 0.736 0.741 0.787 0.823 0.838 0.867 0.950 1.020 1.112 1.201 1.319 1.386 1.702
If the system follows Young’s rule, the mixing in terms of density can be considered as ideal. For comparison, the calculated results by eq 8 are shown in Table 3 combined with the ARD and rmsd. It can be seen that this ideal mixing model results in average absolute deviations between experimental and calculated densities within 0.1 %. The results indicate that Young’s rule is suitable and reliable to extrapolate densities of the aqueous (NH3OH)2SO4 + (NH4)2SO4 system. Similar results have been obtained for the aqueous H2SO4 + FeSO411 and NaOH + NaAl(OH)412,14,15 systems. Viscosities for (NH3OH)2SO4 + H2O Solutions. The viscosities of aqueous (NH3OH)2SO4 + H2O solutions were measured at T ) (293.15, 303.15, and 308.15) K. The experimental viscosities are listed in Table 4. For electrolyte solution, the extended Jones-Dole equation9 in terms of ionic strength, I, was used to express the relation between the viscosity and electrolyte concentration.
ηr )
(8) where F(IT) is the density of the ternary mixture using total ionic strength, IT, in molality units, I(NH3OH)2SO4 and I(NH4)2SO4 are the ionic strength contributions of (NH3OH)2SO4 and (NH4)2SO4 in solution, respectively, and F(NH3OH)2SO4(IT) and F(NH4)2SO4(IT) are the densities of the binary solutions at total ionic strength, IT, respectively. In the present case, the density of aqueous (NH4)2SO4 solution, F(NH4)2SO4(IT), is calculated from the equation and parameters provided by Tans.13
0.0611 0.1418 0.2107 0.6944 1.1127 1.2942 1.6002 2.4206 3.1308 3.9709 4.7676 5.7435 6.2645 8.5406
η ) 1 + AI 1/2 + BI + CI 2 η0
(9)
When a linear dependence of temperature is used, the above equation can be written as
ηr )
η ) 1 + (A1 + A2T)I 1/2 + (B1 + B2T)I + (C1 + η0 C2T)I 2 (10)
In the above equations, η0 is the viscosity of pure water at the selected temperature and Ai, Bi, and Ci are temperature-
Journal of Chemical & Engineering Data, Vol. 54, No. 7, 2009 2031 Table 5. Coefficients of Equation 10 for Viscosity Fits of Hydroxylamine Sulfate in Water A1
A2
B1
B2
C1
C2
100 ARD
rmsd
-0.115593
0.000435
0.060594
0.000148
0.025103
-0.000064
0.15
0.003
independent constants, which are estimated under 0.95 confidence intervals from fitting of the experimental data. The values of these parameters are reported in Table 5 combined with the AAD and rmsd of calculation. These results show that eq 10 can be satisfactorily used for correlation of the experimental viscosity data over a wide range of concentrations. In the Jones-Dole equation, the viscosity B coefficient is a valuable tool to provide information concerning the solvation of solutes and their effects on the structure of the solvent in the near environment of solute molecules or ions.16,17 Table 5 indicates that values of the B (calculated by B1 and B2) of hydroxylamine sulfate in water at different temperatures are positive, suggesting the presence of strong ion-solvent interactions.18 The result also indicates that these interactions are strengthened with a rise in temperature. A number of studies17,18 report that dB/dT, B2, is a better criterion for determining the structure-making/breaking nature of any solute rather than the B coefficient. The positive B2 value indicates the structurebreaking tendency of hydroxylamine sulfate in the studied solvent systems. Viscosities for (NH3OH)2SO4 + (NH4)2SO4 + H2O Solutions. The viscosities of ternary mixtures of (NH3OH)2SO4 + (NH4)2SO4 + H2O are shown in Table 6. The results show that the replacing of NH4+ by NH3OH+ at the same ionic strength makes the solution become more viscous. This may be attributed to the association of the hydroxylamine ion to form dimers and trimers.19 The hydroxylamine ion has a hydroxyl group with self-association and H-bonding properties. The network of selfassociation forces binds the hydroxylamine ion together with itself or water molecules. Extra energy is required to break this interaction. The ideal mixing rule of Young10 has also been used to predict the experimental viscosity data for the studied ternary mixture. It has been shown that Young’s rule expresses the properties of a ternary mixture as the ionic strength-weighted sum of relevant binary solution properties. Thus, Young’s rule for viscosity of an aqueous hydroxylamine sulfate + ammonium sulfate mixture, η(IT), can be written as
η(IT) )
I(NH3OH)2SO4 IT
η(NH3OH)2SO4(IT) +
I(NH4)2SO4 IT
η(NH4)2SO4(IT) (11)
where IT is the toal ionic strength in molality units, I(NH3OH)2SO4 and I(NH4)2SO4are the ionic strength contributions of (NH3OH)2SO4 and (NH4)2SO4 in solution F(NH3OH)2SO4(IT), respectively, and η(NH3OH)2SO4(IT) and η(NH4)2SO4(IT) are the viscosities of the binary solutions at total ionic strength, IT, respectively. The viscosity of the aqueous (NH4)2SO4 solution was calculated from the literature,20 which was determined over a wide range of concentrations and temperatures. The calculated values are listed in Table 6 for comparison. The obtained standard deviations for η at T ) (298.15, 303.15, and 308.15) K are 0.015, 0.012, and 0.023, respectively. The average relative deviation between the experimental and calculated viscosity data is less than 1 %. These results show that the mixing in terms of viscosity can also be considered as ideal and eq 11 can be satisfactorily used for the prediction of ternary viscosities for the title system.
Table 6. Viscosities of Aqueous Hydroxylamine Sulfate + Ammonium Sulfate Mixture at Different Temperatures η/mPa · s +
+
m(NH3OH ) m(NH4 ) T ) 298.15 K T ) 303.15 K T ) 308.15 K mol · kg-1
mol · kg-1
exp
calcda
exp
calcda
exp
calcda
0.0000 0.5325 1.0693 1.3383 1.6073 2.1536 0.2535 1.0792 1.7861 3.1504 2.4698 1.7931 0.8918 1.2916 0.6971
9.3100 8.7906 8.2826 8.0292 7.7681 7.2655 7.3433 6.5382 5.8686 0.2265 0.8970 1.5684 2.4525 0.8607 1.4479 rmsd 100 ARD
2.114 2.219 2.291 2.290 2.319 2.410 1.858 1.929 2.029 1.504 1.442 1.387 1.304 1.200 1.150
2.109 2.181 2.256 2.294 2.331 2.411 1.847 1.948 2.038 1.510 1.448 1.387 1.304 1.199 1.149 0.015 0.49
1.944 2.011 2.051 2.071 2.090 2.157 1.679 1.758 1.825 1.350 1.295 1.248 1.179 1.081 1.041
1.938 1.995 2.055 2.085 2.114 2.178 1.689 1.770 1.843 1.356 1.303 1.251 1.181 1.081 1.038 0.012 0.55
1.785 1.824 1.866 1.879 1.887 1.939 1.531 1.598 1.665 1.223 1.169 1.138 1.069 0.981 0.942
1.781 1.831 1.884 1.911 1.937 1.994 1.547 1.619 1.683 1.222 1.177 1.133 1.073 0.978 0.942 0.023 0.94
a
Calculated by eq 11.
Conclusions The densities and viscosities of the industrially important (NH4)2SO4 + H2O and (NH3OH)2SO4 + (NH4)2SO4 + H2O systems are presented over a wide range of concentrations and temperatures. The Masson-type equation and Jones-Dole equation in terms of ionic strength were used to correlate the experimental binary mixture data. This work demonstrates that the ideal mixing rule of Young in terms of ionic strength can be successfully applied to predict the properties of the ternary (NH3OH)2SO4 + (NH4)2SO4 + H2O system.
Literature Cited (1) Ritz, J.; Fuchs, H.; Perryman, H. G. Hydroxylamine. In Ullmann’s Encyclopedia of Industrial Chemistry, 6th ed.; Wiley: Chichester, 2000. (2) Koff, F. W.; Slfnlades, S.; Tunlck, A. A. Ion-Exchange Process for the Recovery of Hydroxylamine from Raschig Synthesis Mixtures. Ind. Eng. Chem. Process Des. DeV. 1982, 21, 204–210. (3) Slfnlades, S.; Tunlck, A. A.; Koff, F. W. Liquid Ion-Exchange Process for the Recovery of Hydroxylamine from Raschig Synthesis Mixtures. Ind. Eng. Chem. Process Des. DeV. 1982, 21, 211–216. (4) He, C. H.; Liu J. Q.; Zhu M. Q.; Xie F. Y.; He C. Y. Method for Preparation of Solid Hydroxylamine Chloride. C.N. Patent No. 200510060291.4, 2007 (in Chinese). (5) Fang, S.; Zhao, C. X.; He, C. H. Densities and Viscosities of Binary Mixtures of Tri-n-butyl Phosphate + Cyclohexane + n-Heptane at T ) (288.15, 293.15, 298.15, 303.15, and 308.15) K. J. Chem. Eng. Data 2008, 53, 2244–2246. (6) Fang, S.; Zhao, C. X.; He, C. H.; Liu, J. Q.; Sun, J. H. Densities and Viscosities of Binary Mixtures of Tris(2-ethylhexyl) Phosphate + Cyclohexane or n-Hexane at T ) (293.15, 298.15, and 303.15) K and p ) 0.1 MPa. J. Chem. Eng. Data 2008, 53, 2718–2720. (7) Skvortsov, G. A.; Skvortsov, A. B.; Klyushneva, M. N. Physicochemical properties of aqueous solutions of hydroxylamine sulfate. I. Zh. Prikl. Khim. 1971, 44, 1518–1521. (8) Sohnel, O.; Novotny, P. Densities of Aqueous Solutions of Inorganic Substances; Elservier: New York, 1985. (9) Jones, G.; Dole, M. The Viscosity of Aqueous Solutions of Strong Electrolytes with Special Reference to Barium Chloride. J. Am. Chem. Soc. 1929, 51, 2950–2964. (10) Young, T. F.; Smith, M. B. Thermodynamic Properties of Mixtures of Electrolytes in Aqueous Solutions. J. Phys. Chem. 1954, 58, 716– 724. (11) Konigsberger, E.; Konigsberger, L. C.; Szilagyi, I.; May, P. M. Measurement and Prediction of Physicochemical Properties of Liquors
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(16)
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Relevant to the Sulfate Process for Titania Production. 1. Densities in the TiOSO4+ FeSO4 + H2SO4 + H2O System. J. Chem. Eng. Data 2009, 54, 520–525. Sipos, P.; Stanley, A.; Bevis, S.; Hefter, G.; May, P. M. Viscosities and Densities of Concentrated Aqueous NaOH/NaAl(OH)4 Mixtures at 25 °C. J. Chem. Eng. Data 2001, 46, 657–661. Tans, A. M. P. A New Type of Nomogram. Aqueous Ammonium Sulfate Solutions. Ind. Eng. Chem. 1958, 50, 971–972. Magalhaes, M. C. F.; Konigsberger, E.; May, P. M.; Hefter, G. Heat Capacities of Concentrated Aqueous Alkaline Aluminate Solutions at 25 °C. J. Chem. Eng. Data 2002, 47, 960–963. Konigsberger, E.; Bevis, S.; Hefter, G.; May, P. M. Comprehensive Model of Synthetic Bayer Liquors. Part 2. Densities of Alkaline Aluminate Solutions to 90 °C. J. Chem. Eng. Data 2005, 50, 1270– 1276. Rajwade, B. R. P.; Pande, R. Partial Molar Volumes and Viscosity B-Coefficient of N-Phenylbenzohydroxamic Acid in Dimethylsulfoxide at Different Temperatures. J. Chem. Eng. Data 2008, 53, 1458–1461.
(17) Sharma, T. S.; Ahluwalla, J. C. Experimental Studies on the Structures of Aqueous Solutions of Hydrophobic Solutes. ReV. Chem. Soc. 1973, 2, 203–232. (18) Dakua, V. K.; Biswajit, S.; Mahendra, N. R. Ion-solvent and ion-ion interactions of sodium molybdate salt in aqueous binary mixtures of 1,4dioxane at different temperatures. Phys. Chem. Liq. 2007, 45, 549–560. (19) Morris, K. B.; Nelson, J. J.; Nizko, H. S. Electrical Conductance, Density, and Viscosity of Aqueous Solutions of Hydroxylammonium Chloride at 25 °C. J. Chem. Eng. Data 1965, 10, 120–121. (20) Guo, C. Y. The properties of aqueous electrolyte solutions. Master’s Thesis, Providence University, 2007 (in Chinese). Received for review November 15, 2008. Accepted March 27, 2009. This work was supported by the National Key Technology R&D Program of China (Project No. 2007BAI34B07) and the Commission of Science and Technology of Zhejiang Province (Project No. 2008C01006-1).
JE8008649
